Ions and their Hydration

When sulfur(VI) oxide dissolves in water, a solution containing hydrated protons and hydrated hydrogen sulfate(VI) ions [HSOT(aq)] is produced, with some of the latter ions dissociating to give more hydrated protons and hydrated sulfate(VI) ions  

In general, ionic compounds act as true solutes in that evaporation of the solvent yields the original compounds. Evaporation of a solution of HC1 will give gaseous HCI, but evaporation of the S03 solution gives sulfuric(Vl) acid, H2SO4. 

These examples are sufficient to form the basis of the discussion of the nature of cations and anions in aqueous solutions. In aqueous solution, ions are stabilized by their interaction with the solvent they become h drated and this slate is indicated in equations by the (a ) symbolism. This is a broad generalization and is amplified by further discussions in the next sections and in the remainder of the book. 

A simple calculation reveals the limits of the numbers of water molecules that may be associated with an ion in a standard solution. A l mol dm-3 aqueous solution of sodium chloride has a density of 1038 kg m-3 at 25 °C, so 1 dm3 of such a solution has a mass of 1038 g. One mole of the salt has a mass of 58.44 g, so the water in the litre of solution has a mass of 1038-58.44 = 979.56 g. This amount of water contains 979.56/18.015 = 54.4 moles of the liquid. The molar ratio of water molecules to ions in the 1 mol dm-3 aqueous solution of Na h(aq) and Cl (aq) ions is therefore 54.4/2 = 27.2, assuming that the water molecules are shared equally between the cations and anions. This represents the theoretical upper limit of hydration of any one ion in a standard solution of 1 mol dm-3 concentration. The limit may be exceeded in more dilute solutions, but that depends upon the operation of forces over a relatively long range. Certainly, in more concentrated solutions, the limits of hydration of ions become more restricted as fewer water molecules are available to share out between the cation and anion assembly. 

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TABLE 3 Charge Numbers and Radii of Ions and Their Hydration Numbers and Hydrated Radii in NB at 25° C... 

This paper starts with a brief description of the Golden ratio and the ( )-based crystal ionic radii and is then followed by the (()-based aqueous ionic radii and hydration lengths. The role of in the sizes of the ions in the crystal and in aqueous solutions and their hydration bonds with water can be seen in Fig. 12.3 for Na" and Cl" ions (used as the examples). 

The last example of a sequential approach is from Sanov (excerpt 130). A series of increasingly complex experiments is proposed to study the photochemistry of 02, and OCS . Sanov begins with the easier diatomic anions (02 and which will serve as prototypes for subsequent experiments. Next, he will study a larger, polyatomic anion (OCS ) and its cluster ions, 0CS (H20)]j. In the future, he will study even larger dimers and trimers (OCS)n (n > 2) and their hydrated counterparts. 

The study of coordination compounds of the lanthanides dates in any practical sense from around 1950, the period when ion-exchange methods were successfully applied to the problem of the separation of the individual lanthanides,131-133 a problem which had existed since 1794 when J. Gadolin prepared mixed rare earths from gadolinite, a lanthanide iron beryllium silicate. Until 1950, separation of the pure lanthanides had depended on tedious and inefficient multiple crystallizations or precipitations, which effectively prevented research on the chemical properties of the individual elements through lack of availability. However, well before 1950, many principal features of lanthanide chemistry were clearly recognized, such as the predominant trivalent state with some examples of divalency and tetravalency, ready formation of hydrated ions and their oxy salts, formation of complex halides,134 and the line-like nature of lanthanide spectra.135... 

The evidence reviewed here is consistent with the idea that the condensed conformation of secretory products during storage in the cell, and their hydrated conformation upon release from the cell, reflect the corresponding condensed and decondensed phases of a polymer gel. Product release in exocytosis would result from a polymer gel phase transition that is probably triggered by a polycation Zl +2/Na+ ion exchange via the secretory pore. 

The extent of the accumulation of such intermediates depends on their rates of the formation and those of the ensuing decomposition (or dissolution) reactions. If the latter are not high, the total density of such surface intermediates becomes so high that an appreciable surface potential At) is created by the electric double layer formed by the charge unbalance of these intermediates, as well as by the approach of counter ions and the hydration around these intermediates. The experimentally obtained difference between Ug(dark) and Ug (ill.) can be attributed to this AiJj. 

The relative reactivity of the different mineral phases of cement with water is usually given as C A>C S>C S>C AF. Aluminate phases and their hydration products therefore play an important role in the early hydration process. Because of the high reactivity of calcium aluminate, the aluminate hydration reaction is carried out in the presence of sulfate ions. The latter provide control of the reaction rate through the formation of mixed aluminum sulfate products (ettringite and monosulfoaluminate) Calcium sulfate which is added to the cement clinker hence controls the properties of the aluminate hydration products. Sulfates thus play a crucial role in cement hydration and the influence of chemical admixtures on any process where sulfates are involved may be expected to be significant [127],... 

Ruckenstein and Schiby derived4 an expression for the electrochemical potential, which accounted for the hydration of ions and their finite volume. The modified Poisson-Boltzmann equation thus obtained was used to calculate the force between charged surfaces immersed in an electrolyte. It was shown that at low separation distances and high surface charges, the modified equation predicts an additional repulsion in excess to the traditional double layer theory of Deijaguin—Landau—Verwey—Overbeek. 

Na MAS NMR has also been used to study the structure and sodium environment in amorphous sodium aluminosilicate geopolymers, showing that the charge-balancing Na" " is present in a highly hydrated form (Barbosa et al. 2000). When the geopolymer is heated to > 1200°C, the sodium ions lose their hydration water, as evidenced by a shift in the position of the Na resonance from- 5.5 ppm to — 19 ppm, but the amorphous nature of the material is retained. 

Figure 2.3. Surface complexation phenomena in the retention or desorption of metals from mineral surfaces. Nonspecific (exchangeable) adsorption consists of electrostatic bonds only and the ions retain their hydration sphere (outer-sphere complexes) specific (nonexchangeable) adsorption requires removal of the hydration sphere (inner-sphere complexes). Alkali and alkaline earth metals tend to form outer-sphere complexes, hence their tendency to be loosely bound and readily exchangeable with other ions in solution. Transition metals tend to form inner-sphere complexes, which are more strongly bound and less exchangeable (Cotter-Howells and Paterson, 2000). Representation of (a) an outer-sphere complex, (b) an inner-sphere complex, and (c) a solution complex (see also Figure 2.2). The solid substrate is textured with the solution above this. Unlabeled spheres represent oxygen atoms, and the spheres labeled M represent metals in the substrate or in solution. Smaller shaded spheres labeled H are hydrogen atoms. (Adapted from Brown et al., 1999 Cotter-Howells and Paterson, 2000.)...

A specific free radical can be produced from a precursor molecule either in an initiation step or a propagation step in which a reagent radical reacts with the precursor. Initiation requires either removal or addition of an electron or homolysis. Chemically this can be done in a number of ways, by using one-electron oxidants or reductants or by inducing homolysis in some way examples of these types of reactions include autoxidation [84-86], photochemical oxidation and reduction [87-90], and oxidation and reduction by metal ions and their complexes [91-93], In propagation reactions, the reagent radical might be the hydroxyl radical, the hydrated electron, or any other suitably reactive species that will interact with the precursor molecule in the desired manner. We will consider initiation reactions first. 

Finally, considerable progress has been made with regard to gas-phase conformational analysis of monosaccharides and their hydrated complexes by comparison of their IR spectra with that predicted by ab initio calculations.76 78 Many of the tools are therefore available to place a sugar in the gas phase, cleave the glycosidic bond, isolate the oxocarbenium ion, and then have the opportunity to study the kinetics for nucleophilic capture by direct kinetic techniques. Experiments such as competitive KIEs for capture of the oxocarbenium ion also seem possible. Such work may allow direct experimental determination of oxocarbenium ion lifetimes, barriers for capture, and transition state structures. Differences observed between gas-phase results and those in solution may reveal the role that solvent plays in the reaction. 

The heat of adsorption of water on a metal surface is needed in calculations explaining ion adsorption on metals. A major contribution to the free energy of adsorption of an ion is the enthalpy lost by displacement of adsorbed water and by the reorientation of neighboring adsorbed water molecules in the hydration sphere when they have rotated their dipoles due to the influence of the adsorbed ion. Other contributions depend on changes in total hydration of the ion. This is a result of the angular relationship of the water and the ion and their respective distances from the metal surface. 

In his early study covering a large number of halides and oxyhalides of thorium and their hydrates, Chauvenet [1911CHA] reported a value of -293.67 kJ moT for the enthalpy of solution of ThBr4(cr) in ca. 16000 H2O, probably at 288 K. Use, as discussed in Appendix A, of the values adopted by this review (Table VII-15) for the hydrolysis of the thorium ion and of that for the formation of the first thorium bromide complex (Section VIII.2.2) leads to a dissolution reaction that can be written as ... 

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